energy and phases

Energy and Phases

Potential EnergyPotential Energy - stored energy - stored energy (stored in bonds, height)(stored in bonds, height)

Kinetic EnergyKinetic Energy - energy of motion, - energy of motion, associated with heatassociated with heat

Forms of EnergyForms of Energy

Light - light waves, electromagnetic Light - light waves, electromagnetic radiationradiation

ElectricalElectrical ChemicalChemical HeatHeat Mechanical - moving parts, machinesMechanical - moving parts, machines Atomic/Nuclear - changes in mass of Atomic/Nuclear - changes in mass of

atomatom

Conservation of Energy

Energy can be converted from one form to another but never destroyed.

The total amount of energy is always constant.

Exothermic Reactions

• Energy (heat) exits• Releases energy (heat) when new products

are formed• Potential Energy of the reactants is greater

than the potential energy of the products• Surroundings feel warm because heat was

released• Heat is a product

Ex: AB A + B + heat

Endothermic Reactions

• Endothermic – energy (heat) goes in• Reaction absorbs heat from the

environment• Potential energy of the reactants is less

than the potential energy of the products• Surroundings feel cold because heat was

absorbed from the surrounding• Heat is a reactant• Ex: A + B + heat AB

Activation Energy

• The energy needed to break the bonds (to get the reaction started)

• Energy to get over the “hill”, difference between your starting point and the top of the hill

• All reactions need activation energy

Heat of Reaction (H)

• Heat absorbed/released by a reactionH = Hp - HR

• If H is negative, Exothermic

• If H is positive, Endothermic

Table I

1. Is the reaction, C(s) + O2(g) → CO2(g), exothermic or endothermic?

2. What is the value of ΔH for the reaction, CO2(g) → C(s) + O2(g)

Catalyst

• Lowers the activation energy

• Speeds up the reaction

Heat and Temperature

• Heat and temperature are not the same– Heat – measure of the energy transferred from one

substance to another– Temperature – measure of the average kinetic

energy of a substance’s particles

• The faster the particles move (more KE), the higher the temperature

• Heat flows from high to lower until an equilibrium is established (Hot Cold)

• Calorimeter – measures the heat given off by a reaction

Heat/Temp Examples

1. If two systems at different temperatures have contact with each other, heat will flow from the system ata. 20oC to a system at 303Kb. 30oC to a system at 313Kc. 40oC to a system at 293Kd. 50oC to a system at 333K

2. Which is not a form of energy? a. Light b. Temperaturec. Heat d. Motion

3. Which has the most kinetic energy? a. 10.0 g of H2O at 70oC b. 10.0 g of H2O at 5oCc. 25.0 g of H2O at 60oC d. 25.0 g of H2O at 10oC

Solids

• Definite shape• Definite volume• Very high attractive forces between

molecules• Neighboring particles are very

close together• Crystalline structure

• Kinetic Energy – solids have KE– Particles are constantly vibrating

(around their positions in the crystal)

– Positions do not change in relation to the other particles in the crystal

– At absolute zero (O K) all movement stops (theoretically)

Melting Point As energy is added to the solid, KE of the

particles increases until they have sufficient energy to overcome the forces holding them in the crystal, the substance begins to melt

Temperature where the solid and liquid phase exist in equilibrium

Heat of Fusion Heat needed to melt Amount of heat needed to convert a unit mass

of a substance from a solid to liquid at a constant temperature

q = mHf

q = heat (J)m = mass (g)Hf = heat of fusion (J/g)Hf for water = 334 J/g

Hf Examples

1. How much energy is needed to change 75g of ice at OoC to water at the same temperature?

2. 11,000J of heat are released as a sample of water at OoC freezes. Calculate the mass of the sample.

Melting/Freezing Melting is an endothermic process, the H2O

absorbs 334J/g Freezing is the reverse process, so it is

exothermic, the H2O releases 334J/g

Sublimation

• Solid phase gas phase (skip liquid)

• Endothermic

Example: dry ice CO2(s) CO2(g)

Iodine I2(s) I2(g)

Naphthalene (moth balls)

Liquids

• Definite Volume• Take the shape of the

container• High attractive forces

between molecules (but not as high as those found in solids)

Evaporation (Vaporization, Boiling)

• Change from liquid to gas, endothermic

• As temperature increase, rate of evaporation increases

• Increased temperature, more KE, easier to overcome intermolecular forces and to break them

• Condensation - change from gas to liquid, exothermic

Vapor Pressure

• Pressure of a gas on a liquid• In a closed system (sealed container) the

vapor evaporating from the liquid exerts pressure on the liquid

• Vapor Pressure increases as the temperature of the liquid increases

• It has specific values for each substance at any given temperature

• Table H

Boiling Point

• Temperature where vapor pressure = atmospheric pressure

• Water boils at 100oC at 101.3kPa (1atm)– Pressures below 101.3kPa (high elevations),

water boils below 100oC– Pressures greater than 101.3kPa (below sea

level), water boils above 100oC

Table H Examples

1. What is the vapor pressure of water at 105oC?

2. If the pressure is 30kPa, what temperature will water boil at?

3. What pressure is needed for ethanol to boil at 50oC?

4. Which liquid on Table H has the strongest intermolecular forces?

Heat of Vaporization

• Heat to vaporize

• Amount of heat needed to convert a unit mass of a substance from liquid to vapor at a constant temperature

• Does not increase KE, so temperature does not change

Heat of Vaporization

• q = mHv q = heat (J)

m = mass (g)

Hv = heat of vaporization

Hv for water = 2260 J/g

Examples:

1. How much heat must be supplied to evaporate 50.g of H2O at 100oC?

2. 12,750 J of heat are used to boil a sample of water. Calculate the mass of the sample.

Boiling and Condensation

• Boiling is an endothermic process, the H2O absorbs 2260J/g

• Condensation is the reverse process, so it is exothermic, the H2O releases 2260J/g

Deposition

• Reverse of sublimation

• Direct change from the gas to the solid phase, skip liquid

• Exothermic

• Example: Frost

Phase Diagrams / Heating and Cooling Curves

• When a sample of matter is heated its temperature usually increases.

Increase in KE causes an increase in temp.

• Sometimes matter can gain or lose heat without changing temperature.

What is going on at this point?

Points to remember:

• When heat is used to increase the speed of particles, the temperature increases – at this point the KE is changing

• When heat does not cause a change in temperature, it is being used to change phases (phase changes occur at the flat parts of the graph) – at this point PE is changing

• Melting Point - point when the solid begins to melt, both the solid and liquid phases are present

• Boiling Point - point when the liquid begins to boil, both the liquid and gas phases are present

Rate of Heating

• Total amount of heat absorbed = time x rate of heating

q= t x rate

Example:

A solid is heated at a constant rate of 150 J/min for 3 minutes. How much heat is absorbed?

Heating Curve

Cooling Curve

Change in Temperature Calcs

• q = mcT

• q = heat (Joules)

• m = mass

• C = specific heat capacity

(for water; c = 4.18 J/goC)T = change in temperature

Examples

1. How many joules does it take to raise the temperature of a 5.0g sample of water by 20.oC?

2. A 1000. gram mass of water in a calorimeter has its temperature raised 5.0oC. How much heat energy was transferred to the water?